Chapter Seven. Chemical Bonding and Molecular Structure. Chapter Seven Slide 1 of 98

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1 Chapter Seven Chemical Bonding and Molecular Structure Chapter Seven Slide 1 of 98

2 Chemical Bonds: A Preview Forces called chemical bonds hold atoms together in molecules and keep ions in place in solid ionic compounds. Chemical bonds are electrical forces; they reflect a balance in the forces of attraction and repulsion between electrically charged particles. Through appropriate measurements, scientists can determine the internuclear distances that correspond to the lowest energy states of molecules. Quantum mechanical calculations can then be used to develop a theoretical model that fits the experimental measurements. Chapter Seven Slide 2 of 98

3 Electrostatic Attractions and Repulsions Chapter Seven Slide 3 of 98

4 Energy of Interaction AB A B A B A B Chapter Seven Slide 4 of 98

5 Electronegativity Electronegativity (EN, expressed as χ), is a measure of the ability of an atom to attract bonding electrons to itself when the atom is in a molecule. Mulliken s EN Absolute EN, χ = (IE + EA)/2 Pauling s EN define χ H = 2.1 χ = χ A - χ B = [D AB (kj)/96.49] 1/2 = [D AB (kcal)/23.06] 1/2 D AB = D(A-B) - 1/2 [D(A-A) + D(B-B)] Bond dissociation energy of A-B Chapter Seven Slide 5 of 98

6 Pauling s Electronegativities Chapter Seven Slide 6 of 98

7 Chapter Seven Slide 7 of 98

8 Electronegativity Difference and Bond Type Two identical atoms have the same electronegativity and share a bonding electron pair equally. This is called a non-polar covalent bond. (For χ A ~ χ B, or lχ A - χ B l< 0.3) In covalent bonds between atoms with somewhat larger electronegativity differences (0.3 < lχ A - χ B l < 1.8), electron pairs are shared unequally. The electrons are drawn closer to the atom of higher electronegativity, and the bond is called a polar covalent bond. With still larger differences in electronegativity (lχ A - χ B l> 1.8), electrons may be completely transferred from metal to nonmetal atoms to form ionic bonds.

9 Electronegativity and Bond Type Chapter Seven Slide 9 of 98

10 Ionic Bonds and Ionic Crystals When atoms lose or gain electrons, they acquire a noble gas configuration, but do not become noble gases. Because the two ions formed in a reaction between a metal and a non-metal have opposite charges, they are strongly attracted to one another and form an ion pair. The net attractive electrostatic forces that hold the cations and anions together are ionic bonds. The highly ordered solid collection of ions is called an ionic crystal. Chapter Seven Slide 10 of 98

11 Interionic Forces of Attraction - Coulombic Force E = (Z + Z - )/4πεr

12 Formation of a Crystal of Sodium Chloride Chapter Seven Slide 12 of 98

13 Unit Cell of Sodium Chloride

14 Energy Changes in Ionic Compound Formation The enthalpy of formation of the ionic compound is more important than first ionization energy and electron affinity. The overall enthalpy change can be calculated using a step-wise procedure called the Born-Haber cycle. The sum of the enthalpy change values for the individual steps is in accordance with Hess s Law. The large negative value of the lattice energy is the major factor that makes ionic compound formation an energetically favorable process. Chapter Seven Slide 14 of 98

15 A Born-Haber Cycle Example IE(Na) - EA(Cl) 1/2D(Cl-Cl) Η sublimation (Na) H 5 = - U U : lattice energy Chapter Seven Slide 15 of 98

16 A Born-Haber Cycle Example H 0 f (NaCl) = H 1 + H 2 + H 3 + H 4 + H 5 = Η sublimation (Na) + ½ D(Cl-Cl) + IE(Na) - EA(Cl) - U lattice energy, U = - Η 0 f (NaCl) + Η sublimation (Na) + ½ D(Cl-Cl) + IE(Na) - EA(Cl) = ( ) kj/mol = +787 kj/mol Chapter Seven Slide 16 of 98

17 The Lewis Theory of Chemical Bonding Electrons, particularly valence electrons, play a fundamental role in chemical bonding. When metals and non-metals combine, valence electrons usually are transferred from the metal to the non-metal atoms giving rise to ionic bonds. In combinations involving only non-metals, one or more pairs of valence electrons are shared between the bonded atoms producing covalent bonds. In losing, gaining, or sharing electrons to form chemical bonds, atoms tend to acquire the electron configurations of noble gases. Chapter Seven Slide 17 of 98

18 Lewis Symbols In a Lewis symbol, the chemical symbol for the element represents the nucleus and core electrons of the atom, and dots around the symbol represent the valence electrons. Chapter Seven Slide 18 of 98

19 Using Lewis Symbols to Represent Ionic Bonding Instead of using complete electron configurations to represent the loss and gain of electrons, Lewis symbols can be used for ionic bond. Na + Cl: Na 1+ :Cl: 1- Chapter Seven Slide 19 of 98

20 Lewis Structures Of Simple Molecules A Lewis structure is a combination of Lewis symbols that represents the formation of covalent bonds between atoms. In most cases, a Lewis structure shows the bonded atoms with the electron configuration of a noble gas; that is, the atoms obey the octet rule. (H obeys the duet rule.) Chapter Seven Slide 20 of 98

21 Lewis Structures (continued) The shared pairs of electrons in a molecule are called bonding pairs. In common practice, the bonding pair is represented by a dash (-). The other electron pairs, which are not shared, are called non-bonding pairs, or lone pairs. or Chapter Seven Slide 21 of 98

22 2 nd period elements follow octet rule The non-metals of the second period (except boron?) tend to form a number of covalent bonds equal to eight minus the group number. Chapter Seven Slide 22 of 98

23 Coordinate Covalent Bonds In some cases, one atom provides both electrons of the shared pair to form a bond called a coordinate covalent bond (or dative bond). For example: hydronium ions H 3 O + F 3 B-NH 3 adduct Chapter Seven Slide 23 of 98

24 Multiple Covalent Bonds The covalent bond in which one pair of electrons is shared is called a single bond. Multiple bonds can also form: Double bonds have two shared pairs of electrons. Triple bonds have three shared pairs of electrons. A double bond is represented by two dashes (=). A triple bond is represented by three dashes ( ). Chapter Seven Slide 24 of 98

25 Polar and Non-polar Covalent Bonds Chapter Seven Slide 25 of 98

26 Intermolecular interaction Chapter Seven Slide 26 of 98

27 Formal Charge Formal charge = # of valence electrons - [# of lone pair electrons + ½ (bonding electrons)] Usually, the most plausible Lewis structure is one with no formal charges. When formal charges are required, they should be as small as possible. Negative formal charges should appear on the most electronegative atoms. Adjacent atoms in a structure should not carry formal charges of the same sign. Chapter Seven Slide 27 of 98

28 Lewis Structures & Formal Charge CO FC(C) = = -1 (-1) (+1) FC(O) = = +1 CO 2 (0) (0) (0) FC(C) = = 0 FC(O) = = 0 N 2 O resonance structures (-1) (+1) (0) (0) (+1) (-1) (-2) (+1) (+1) (-1) (+2) (-1) Chapter Seven Slide 28 of 98

29 Resonance: Delocalized Bonding Resonance theory states that whenever a molecule or ion can be represented by two or more plausible Lewis structures that differ only in the distribution of electrons, the true structure is a composite, or hybrid, of them. The different plausible structures are called resonance structures. The actual molecule or ion that is a hybrid of the resonance structures is called a resonance hybrid. Electrons that are part of the resonance hybrid are spread out over several atoms and are referred to as being delocalized. Chapter Seven Slide 29 of 98

30 Oxoacids: HNO 3, H 2 SO 4, HClO 4 ClO 4 - (-1) (0) (-1) (-1) (-1) (-1) (+3) (+2) (+1) (-1) (-1) (-1) (0) (-1) (-1) Most possible resonance forms (-1) Most possible resonance forms Chapter Seven Slide 30 of 98

31 Molecules that Don t Follow the Octet Rule Molecules with an odd number of valence electrons have at least one of them unpaired and are called free radicals. Some molecules have incomplete octets. These are usually compounds of Be, B, and Al, generally have some unusual bonding characteristics, and are often quite reactive. Chapter Seven Slide 31 of 98

32 Molecules that Don t Follow the Octet Rule (continued) Some compounds have expanded valence shells, which means that the central atom has more than eight electrons around it. e.g. SF 6, PCl 5 Chapter Seven Slide 32 of 98

33 Molecular Geometry The molecular geometry, or the shape of a molecule is described by the geometric figure formed when the atomic nuclei are imagined to be joined by the appropriate straight lines. Determination of molecular geometry 1. Valence-Shell Electron-Pair Repulsion (VSEPR) - Pairs of valence electrons in bonded atoms repel one another. - The mutual repulsions push electron pairs as far from one another as possible. 2. Ligand Field Stabilization Energy - Mainly for compounds with the central atoms as transition metals Chapter Seven Slide 33 of 98

34 Molecular Geometry of Water Chapter Seven Slide 34 of 98

35 Electron-Group Geometries An electron group is any collection of valence electrons, localized in a region around a central atom, that repels other groups of valence electrons. The mutual repulsions among electron groups lead to an orientation of the groups that are called electrongroup geometry. Chapter Seven Slide 35 of 98

36 Electron-Group Geometries #. of Electron-group Electron Pair Geometry 2 linear 3 trigonal planar 4 tetrahedral 5 Trigonal bipyramidal 6 octahedral Chapter Seven Slide 36 of 98

37 A Balloon Analogy Chapter Seven Slide 37 of 98

38 Geometries of Methane #. of electron group on C = 1/2 (4 + 4) = 4 Electron group geometry: tetrahedral VSEPR notation AX 4 Molecular geometry: tetrahedral Chapter Seven Slide 38 of 98

39 Molecular Geometry of Water #. of electron group on O = 1/2 (6 + 2) = 4 Electron group geometry: tetrahedral VSEPR notation AX 2 E 2 Molecular geometry: bent A: the central atom in a structure X: terminal atoms E: the lone pairs of electrons Chapter Seven Slide 39 of 98

40 Electron Repulsion Lone-pair Lone-pair > Lone-pair Bonding-pair > Bonding-pair Bonding-pair Triple bond > Double bond > Single bond Chapter Seven Slide 40 of 98

41 Molecular Geometry for Iodine Pentafluoride (IF 5 ) #. of electron group on I = 1/2 (7 + 5) = 6 Electron group geometry: octahedral VSEPR notation AX 5 E Molecular geometry: Square pyramidal Chapter Seven Slide 41 of 98

42 Molecular Geometry for Sulfur Tetrafluoride (SF 4 ) #. of electron group on S = 1/2 (6 + 4) = 5 Electron group geometry: trigonal bipyramidal VSEPR notation AX 4 E Molecular geometry: trigonal pyramidal or Seesaw Chapter Seven Slide 42 of 98

geometry: trigonal bipyramidal VSEPR notation AX 3 E 2

43 Molecular Geometry for Iodine Trifluoride (IF 3 ) #. of electron group on I = 1/2 (7 + 3) = 5 Electron group geometry: trigonal bipyramidal VSEPR notation AX 3 E 2 Molecular geometry: triangular or T-shaped Chapter Seven Slide 43 of 98

44 Molecular Geometry for Iodine Tetrafluoride Ion (IF 4- ) #. of electron group on I = 1/2 ( ) = 6 Electron group geometry: octahedral VSEPR notation AX 4 E 2 Molecular geometry: Square planar Chapter Seven Slide 44 of 98

45 Polar Molecules and Dipole Moments A molecule with separate centers of positive and negative charge is called a polar molecule. The dipole moment (µ) of a molecule is the product of the magnitude of the charge (δ) and the distance (d) that separates the centers of positive and negative charge. µ = δ d Chapter Seven Slide 45 of 98

46 Dipole moments are generally expressed in a quantity called a debye. 1 debye (D) = 3.34 x C m = esu charge of electron = x C = 4.80 x esu Calculation of percentage ionic character of HCl Η Cl, µ = 1.08 D, r H-Cl = 127pm δ = µ /d = 1.08 x 3.34 x /(127 x ) = 2.84 x10-20 C percentage ionic character = [2.84 x10-20 / (1.602 x )] x 100% = 17.7% Chapter Seven Slide 46 of 98

47 Dipole-Dipole Interactions d µ 1 Dipole moment r µ = δ d µ 2 E = - 2(µ 1 µ 2 )/4πεr 3 Chapter Seven Slide 47 of 98

48 How Polar Molecules Behave in an Electric Field Chapter Seven Slide 48 of 98

49 Bond Dipoles and Molecular Dipoles All polar covalent bonds have a bond dipole; a separation of positive and negative charge centers in an individual bond. Bond dipoles have both a magnitude and a direction. molecular dipole = vector sum of bond dipoles CO 2, linear with no dipole moment (µ = 0 D) water is bent (bond angle = o ) and µ = 1.84 D. Chapter Seven Slide 49 of 98

50 Molecular Shapes and Dipole Moments Molecules can be predicted to be polar or non-polar based on the following three-step approach: Use electronegativity values to predict bond dipoles. Use the VSEPR method to predict the molecular shape. From the molecular shape, molecular dipole = vector sum of bond dipoles Lone-pair electrons can also make a contribution to dipole moments. e.g. NH 3 µ = 5.0 x C.m NF 3 µ = 0.7 x C.m Chapter Seven Slide 50 of 98

51 Atomic Orbital Overlap Valence Bond (VB) Theory states that a covalent bond is formed by the pairing of two electrons with opposing spins in the region of overlap of atomic orbitals between two atoms. This overlap region has a high electron charge density. In general, the more extensive the overlap between two orbitals, the stronger is the bond between two atoms. The valence bond theory attempts to find the best approximation of optimal orbital overlap for all the bonds in a molecule. Chapter Seven Slide 51 of 98

52 Bonding in H 2 Chapter Seven Slide 52 of 98

53 Bonding In H 2 S Chapter Seven Slide 53 of 98

54 Several Important Points Most of the electrons in a molecule remain in the same orbital locations that they occupied in the separated atoms. Bonding electrons are localized in the region of atomic orbital overlap. For orbitals with directional lobes, maximum overlap occurs when atomic orbitals overlap end to end; that is, a hypothetical line joining the nuclei of the bonded atoms passes through the region of maximum overlap. The molecular geometry depends on the geometric relationships among the atomic orbitals of the central atom that participate in bonding. Chapter Seven Slide 54 of 98

55 Hybridization Of Atomic Orbitals sp 3 Hybridization Scheme Chapter Seven Slide 55 of 98

56 Bonding in Ammonia Chapter Seven Slide 56 of 98

57 The sp 2 Hybridization Scheme Chapter Seven Slide 57 of 98

58 The sp Hybridization Scheme Chapter Seven Slide 58 of 98

59 Hybrid Orbitals Involving d Subshells This hybridization allows for expanded valence shell compounds. A 3s electron can be promoted to a 3d subshell which gives rise to a set of five sp 3 d hybrid orbitals. These molecules have a trigonal bipyramidal molecular geometry. One 3s electron and one 3p electron can be promoted to two 3d subshells which gives rise to a set of six sp 3 d 2 hybrid orbitals. These molecules have an octahedral molecular geometry. Chapter Seven Slide 59 of 98

60 Hybrid Orbitals and Their Geometric Orientations Chapter Seven Slide 60 of 98

61 Hybrid Orbitals and Multiple Covalent Bonds Covalent bonds formed by the end-to-end overlap of orbitals, regardless of orbital type, are called sigma (σ) bonds. All single bonds are sigma bonds. No nodal plane along inter-nuclear axis A bond formed by parallel, or side-by-side, orbital overlap is called a pi (π) bond. One nodal plane along inter-nuclear axis Chapter Seven Slide 61 of 98

62 Descriptions of Ethylene Rotational barrier for double bond Chapter Seven Slide 62 of 98

63 Valence Bond Theory of the Bonding in Acetylene Chapter Seven Slide 63 of 98

64 Geometric Isomerism Geometric isomers are isomers that differ only in the geometric arrangement of certain substituent groups. Two main types of geometric isomers: cis: substituent groups are on the same side trans: substituent groups are on opposite sides Usually formed across double bonds, in cyclic and square planar compounds. Chapter Seven Slide 64 of 98

65 Representative Bond Lengths & Average Bond Energies??? Chapter Seven Slide 65 of 98

66 Some bond energy trends (kj/mol): H-X: H (436) F (567) Cl (432) Br (365) I (298) OH (500) NH 2 (462) H 3 C-X: H (441) F (457) Cl (352) Br (294) I (235) OH (382) NH 2 (365) C-F compounds Teflon PTFE (polytetrafluoroethylene) CFC (chlorofluorocarbons) Chapter Seven Slide 66 of 98

67 Characteristics of Molecular Orbitals Molecular orbitals (MOs) are mathematical equations that describe the regions in a molecule where there is a high probability of finding electrons. Bonding molecular orbitals (σ, π) are at a lower energy level than the separate atomic orbitals and have a high electron probability, or electron charge density. Antibonding molecular orbitals (σ*, π*) are at a higher energy level than the separate atomic orbitals and places a high electron probability away from the region between the bonded atoms. Chapter Seven Slide 67 of 98

68 Molecular Orbitals and Bonding in the H 2 Molecule Chapter Seven Slide 68 of 98

69 The 1s Orbital Ψ (r,θ,φ) = R(r) Υ 0,0 (θ,φ) Υ 0,0 (θ,φ) = 1/2π 1/2 Chapter Seven Slide 69 of 98

70 Constructive Interference + + ψ bonding = φ1s ( A) + φ1s ( B) ψ bonding 2 Chapter Seven Slide 70 of 98

71 Destructive Interference + ψ antibondin g = φ1s ( A) φ1s ( B) - ψ antibondin g 2 Chapter Seven Slide 71 of 98

72 H 2 Bond order = ½ ( # of bonding electrons - # of antibonding electrons ) Electron configuration of H 2 : (σ 1s ) 2 B.O. of H 2 = ½ (2-0) = 1 Bond energy = 435 kj/mol Bond length = 74 pm Chapter Seven Slide 72 of 98

73 H 2 + H 2 - He 2 Chapter Seven Slide 73 of 98

74 Species Electron configuration B.O. Bond energy (kj/mol) Bond length (pm) H 2 (σ 1s ) H 2 + (σ 1s ) 1 ½ H 2 - (σ 1s ) 2 (σ 1s *) 1 ½ He 2 (σ 1s ) 2 (σ 1s *) Chapter Seven Slide 74 of 98

75 Hetero-nuclear Diatomic Molecule Lewis Structure Chapter Seven Slide 75 of 98

76 2nd Period Homo-nuclear Diatomic Molecules Electron configuration of Li 2 : KK(σ 1s ) 2 (σ 1s *) 2 B.O. of Li 2 = ½ (2-0) = 1 Bond length = 267 pm Chapter Seven Slide 76 of 98

77 _ Molecular Orbitals Formed by Combining 2p Atomic Orbitals + _ _ + _ + + _ + _ 1 node inter-nuclear axis _ + 1 node along inter-nuclear axis Chapter Seven Slide 77 of 98

78 Diamagnetic Paramagnetic Chapter Seven Slide 78 of 98

79 Paramagnetic Chapter Seven Slide 79 of 98

80 Paramagnetism of Oxygen Chapter Seven Slide 80 of 98

81 Molecular Orbitals of Homo-nuclear Diatomic Molecules of 2 nd Period Chapter Seven Slide 81 of 98

82 Bonding in Benzene The structure of benzene (C 6 H 6 ), discovered by Michael Faraday in 1825, was not figured out until 1865 by F.A. Kekulé. Kekulé discovered that benzene has a cyclic structure and he proposed that a hydrogen atom was attached to each carbon atom and that alternating single and double bonds joined the carbon atoms together. This kind of structure gives rise to two important resonance hybrids and leads to the idea that all three double bonds are delocalized across all six carbon atoms. Chapter Seven Slide 82 of 98

83 The σ-bonding Framework Chapter Seven Slide 83 of 98

84 The π-molecular Orbitals of Benzene _ E node + + _ node π-m.o. of benzene Chapter Seven Slide 84 of 98 +

85 3 nodes 2 nodes Chapter Seven Slide 85 of 98

86 Aromatic Compounds Many of the first benzene-like compounds discovered had pleasant odors and hence acquired the name aromatic. In modern chemistry, the term aromatic compound simply refers to a substance with a ring structure and with bonding characteristics and properties related to those of benzene. All organic compounds that are not aromatic are called aliphatic compounds. Chapter Seven Slide 86 of 98

87 Some Representative Aromatic Compounds Chapter Seven Slide 87 of 98

88 Chapter Seven Slide 88 of 98

89 Chapter Seven Slide 89 of 98

90 Conjugated Double Bonds E Antibonding Bonding π-m.o. Chapter Seven Slide 90 of 98

91 Band Theory This is a quantum-mechanical treatment of bonding in metals. The spacing between energy levels is so minute in metals that the levels essentially merge into a band. When the band is occupied by valence electrons, it is called a valence band. A partially filled or low lying empty band of energy levels, which is required for electrical conductivity, is a conduction band. Band theory provides a good explanation of metallic luster and metallic colors. Chapter Seven Slide 91 of 98

92 Antibonding bonding Chapter Seven Slide 92 of 98

93 The 2s Band in Lithium Metal Anti-bonding Conduction band e- e- Bonding Valence band Chapter Seven Slide 93 of 98

94 Band Overlap in Magnesium Conduction band Valence band Chapter Seven Slide 94 of 98

95 Band Structure of Insulators and Semiconductors Chapter Seven Slide 95 of 98

96 Chapter Seven Slide 96 of 98

97 Summary The VSEPR method is used to predict the shapes of molecules and polyatomic ions. If all electron-groups are bonding groups, the molecular geometry is the same as the electron-group geometry. A polar covalent bond has separate centers of positive and negative charge, creating a bond dipole. In the valence bond theory, a covalent bond is formed by the overlap of atomic orbitals of the bonded atoms in a region between the atomic nuclei. Hybridized orbitals include sp, sp 2, sp 3, sp 3 d, and sp 3 d 2. Chapter Seven Slide 97 of 98

98 Summary (continued) Unhybridized p orbitals overlap in a side-by-side fashion to form π bonds. Single bonds are all hybridized σ bonds, double bonds have one σ bond and one π bond, and triple bonds have one σ bond and two π bonds. In molecular orbital theory, atomic orbitals of separated atoms are combined into molecular orbitals. The benzene molecule is usually represented by its resonance hybrid. Benzene-like compounds are called aromatic compounds. Chapter Seven Slide 98 of 98

Chapter Nine. Chapter Nine. Chemical Bonds: A Preview. Chemical Bonds. Electrostatic Attractions and Repulsions. Energy of Interaction

Chapter Nine. Chapter Nine. Chemical Bonds: A Preview. Chemical Bonds. Electrostatic Attractions and Repulsions. Energy of Interaction 1 Chemical Bonds: A Preview 2 Chemical Bonds Forces called chemical bonds hold atoms together in molecules and keep ions in place in solid ionic compounds. Chemical bonds are electrostatic forces; they